ap chemistry 4-1 determination of iron by redox titration

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A P CHEMISTRY
Lab 4-1 Determination of Iron by Redox Titration
Pre-Lab Questions - (Must be completed before lab work begins.)
1. Another common analysis for iron is titration of iron (II) with potassium dichromate, K 2Cr2O7.
Divide the following reaction into half-reactions, balance the half-reactions, and combine the
balanced half-reactions into the balanced overall equation for the reaction.
Fe+2 + Cr2O7-2 + H+1  Fe+3 + Cr+3 + H2O
2. If 26.23 mL of potassium permanganate solution is required to titrate 1.041 g of ferrous ammonium
sulfate hexahydrate, FeSO4 (NH4)2SO4 . 6H2O, calculate the molarity of the KMnO4 solution.
2. If a 2.893 g sample of an unknown iron requires 28.45 mL of the permanganate solution described in
Pre-Lab question 3 to reach the endpoint, calculate the percent Fe in the solution.
INTRODUCTION - Oxidation/reduction processes form one of the major classes of chemical
reactions. In redox reactions, electrons are transferred from one species to another. For example, in the
following simple reaction
Zn(s) + Cu+2(aq)  Zn+2(aq) + Cu(s)
electrons are transferred from elemental metallic zinc to aqueous copper (II) ions. This is most easily
seen if the overall reaction is written instead as two half-reactions, one for the oxidation and one for the
reduction:
Zn(s)  Zn +2(aq) + 2eoxidation:
+2
Cu (aq) + 2e  Cu(s)
reduction
In this experiment you will use potassium permanganate, KMnO 4, as the titrant in the analysis of an
unknown sample containing iron. Although this method can easily be applied to the analysis of realistic
iron ores - the native state in which an element is found in nature - for simplicity your unknown sample
will contain only iron in the +2 oxidation state. In acidic solution, potassium permanganate rapidly and
quantitatively oxidizes iron (II) to iron (III), while itself being reduced to manganese (II). The half
reactions for the process are:
MnO4 -1+ 8H+1 + 5e-  Mn+2 + 4H2O reduction
Fe+2  Fe+3 + eoxidation
When these half-reactions are combined to give the overall balanced chemical reaction equation, a
factor of five has to be used with the iron half-reaction so that the number of electrons lost in the overall
oxidation will equal the number of electrons gained in the reduction:
MnO4 + 8H+1 + 5Fe+2  Mn+2 + 4H2O + 5Fe+3
Potassium permanganate has been one of the most commonly used oxidizing agents because it is
extremely powerful, as well as inexpensive and readily available. It does have some drawbacks,
however. Because KMnO4 is such a strong oxidizing agent, it reacts with practically anything that can
be oxidized: this tends to make solutions of KMnO4 difficult to store without decomposition or a change
in concentration. Because of this limitation, it is common to prepare, standardize, and then use KMnO 4
solutions for an analysis all on the same day. It is not possible to prepare directly KMnO4 standard
solutions determinately by mass: solid potassium permanganate cannot be obtained in a completely pure
state due to the high reactivity mentioned above. Rather, potassium permanganate solutions are prepared
to be an approximate concentration, and are then standardized against a known primary standard sample
of the same substance that is to be analyzed in the unknown sample.
Potassium permanganate is especially useful among titrants since it requires no indicator to signal the
endpoint of a titration. Potassium permanganate solutions - even at fairly dilute concentrations - are
intensely colored purple. The product of the permanganate reduction half-reaction, manganese (II), in
dilute solution shows almost no color. Therefore, during a titration using KMnO4, when one drop excess
of potassium permanganate has been added to the sample, the sample will take on a pale red/pink color
(since there are no more sample molecules left to convert the purple MnO 4-1 ions to the colorless Mn+2
ions).
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Lab 4-1 Determination of Iron by Redox Titration
SAFETY:
⇒ Wear safety glasses at all times while in the laboratory.
⇒ Potassium permanganate is a strong oxidizing agent and can be damaging to skin, eyes, and clothing.
Wash after handling.
⇒ Potassium permanganate solutions will stain skin and clothing if spilled.
⇒ Sulfuric acid solutions are damaging to the skin, eyes, and clothing-especially if allowed to
concentrate through evaporation of water. If the sulfuric acid solution is spilled on the skin, wash
immediately and inform the instructor.
⇒ Iron salts may be irritating to the skin. Wash after handling.
Procedure
A. Preparation of Potassium Permanganate Solution
1. Clean out a 1-L glass bottle and its cap. Wash first with soap solution, then rinse with tap water to
remove all the soap. Finally rinse with several small portions of distilled water. Then place
approximately 600 mL of distilled water in the bottle.
2. In a small beaker, weigh out 2.0 ± 0.1 g of potassium permanganate crystals (Caution!). Using a
funnel, transfer the permanganate crystals to the 1-L bottle. Cap the bottle securely and shake the
mixture to dissolve the permanganate. Potassium permanganate is often very slow in dissolving:
continue shaking the solution until it is absolutely certain that all the permanganate crystals have
dissolved completely.
B. Preparation of the Iron Standards
3. Clean out three 250-mL Erlenmeyer flasks with soap and water, giving a final rinse with several
small portions of distilled water. Label the flasks as Samples 1, 2, and 3.
4. The primary standard iron (II) compound to be used for the standardization of the potassium
permanganate solution is the salt ferrous ammonium sulfate hexahydrate (Mohr's salt), FeSO 4
(NH4)2SO4 . 6 H2O, which is almost universally abbreviated as FAS.
5. Place approximately 3 g of primary standard FAS into a small, clean dry beaker. Determine the mass
of the beaker and FAS to at least the nearest milligram (0.001 g), or better, if an analytical balance is
available, make the mass determination to the nearest 0.1 milligram (0.0001 g).
6. Using a funnel, carefully transfer approximately one-third of the FAS in the beaker (approximately 1
g) to the Erlenmeyer flask labeled Sample 1. Distilled water from a wash bottle can be used to rinse
the salt into the flask if the funnel should clog.
7. Weigh the beaker containing the remainder of the FAS. The difference in mass from the previous
weighing represents the mass of FAS transferred to the Sample 1 Erlenmeyer flask.
8. Using a funnel, carefully transfer approximately half the FAS remaining in the beaker (one-third of
the original FAS sample) to the Erlenmeyer flask labeled Sample 2. Again, rinse the FAS from the
funnel into the flask with distilled water from a wash bottle.
9. Weigh the beaker containing the remaining FAS. The difference in mass from the previous weighing
represents the mass of Sample 2 that was transferred to the Erlenmeyer flask.
10. Using a funnel as above, transfer as much as possible of the FAS remaining in the beaker to the
Sample 3 Erlenmeyer flask. Finally weigh the empty beaker and any residual FAS that may not have
been transferred completely. The difference in mass from the previous weighing represents the mass
of Sample 3.
11. Add 25 mL of distilled water to each sample and swirl the flasks to dissolve the FAS.
12. Add 15 mL of 3 M sulfuric acid, H2SO4, to each sample (Caution!). Sulfuric acid is added to the
samples to provide the hydrogen ions, H+1, required for the reduction of the permanganate ion.
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Lab 4-1 Determination of Iron by Redox Titration
C. Standardization of the Potassium Permanganate Solution
13. Clean out a burette with soap and water, and rinse with several portions of tap water, followed by
several small rinsings with distilled water.
14. Rinse the burette with several 4-6 mL portions of the potassium permanganate solution. Tilt and
rotate the burette so that the inside walls are completely rinsed with the permanganate, and allow the
solution to run out of the tip of the burette to rinse the tip also. Discard the rinsings.
15. Finally, with the burette below eye-level, fill the burette to slightly above the zero mark with the
permanganate solution. Allow a few mL of permanganate to run out of the tip of the burette to
remove any air bubbles from the tip.
16. Since potassium permanganate solutions are so intensely colored, it is generally impossible to see the
curved meniscus that the solution surface forms in the burette. In this case, it is acceptable to make
your liquid level readings at the point where the top surface of the permanganate solution comes in
contact with the wall of the burette. Take the initial volume reading, to the nearest 0.02 mL. Record.
17. Place the Erlenmeyer flask containing Sample 1 under the tip of the burette, and begin adding
potassium permanganate to the sample a few mL at a time, swirling the flask after each addition of
permanganate. As the permanganate solution is added to the sample, red streaks may be visible in the
sample until the permanganate has a chance to mix with and react with the iron (II) present.
18. Continue adding permanganate a few mL at a time, with swirling, until the red streaks begin to
become more persistent. At this point, begin adding the permanganate one drop at a time, swirling
the flask constantly to mix. The endpoint is the first appearance of permanent, pale pink color.
19. After the endpoint has been reached, record the final burette reading (to the nearest 0.02 mL).
Calculate the volume of potassium permanganate solution used to titrate FAS Sample 1.
20. Titrate FAS Samples 2 and 3 in a similar manner, recording initial and final volumes of potassium
permanganate solution used to the nearest 0.02 mL.
21. From the mass of each FAS sample, and from the volume of KMnO4 used to titrate the respective
samples, calculate three values for the molarity of your potassium permanganate solution, as well as
the average molarity. If any one of your individual molarities seems out-of-line from the other two,
you should consult with the instructor about performing a fourth standardization titration (using an
additional 1-g sample of FAS).
D. Analysis of the Iron Unknown
22. Your instructor will provide you with an iron unknown, and will inform you of the approximate
sample size to weigh out. Prepare three samples of the iron unknown as you performed in Part B for
the primary standard FAS, being sure to add the sulfuric acid required for the reaction, and being
sure to keep the masses of the samples to the amount indicated by the instructor.
23. Titrate the iron unknown samples as described in Part C above, recording initial and final volumes to
the nearest 0.02 mL. Calculate the volume of permanganate solution required to reach the pale pink
endpoint for each sample.
24. From the volume required to titrate each unknown sample, and the average molarity of the KMnO 4
solution, calculate the number of moles of potassium permanganate required for each iron sample.
25. From the number of moles of potassium permanganate required for titration of each iron sample,
calculate the number of moles of iron present in each unknown sample.
26. From the number of moles of iron present in each sample, and the molar mass of iron, calculate the
mass of iron present in each unknown sample.
27. From the mass of iron present in each sample, and from each sample's mass, calculate the percent
iron in each sample.
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Lab 4-1 Determination of Iron by Redox Titration
Name _________________________________
ANSWER SHEET
PERIOD _____
Results / Observations
Standardization of KMnO4 Solution
Trial 1
Mass of FAS taken
__________
Initial KMnO4 volume
__________
Final KMnO4 volume
__________
Volume KMnO4 used
__________
Moles iron present
__________
Moles KMnO4 present
__________
Molarity of KMnO4 solution
__________
Mean molarity and average deviation
Trial 2
Trial 3
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
____________________
Analysis of Unknown
Trial 1
Trial 2
Mass of unknown taken
__________
Initial KMnO4 volume
__________
Final KMnO4 volume
__________
Volume KMnO4 used
__________
Moles KMnO4 present
__________
Moles iron present
__________
Mass of iron present
__________
Percent of iron present
__________
Mean % iron present and average deviation
Trial 3
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
__________
____________________
Questions
1. Typically, a solid iron unknown for titration is dried in an oven to remove adsorbed water before
analysis (the unknowns used in this experiment were dried before dispensing). How would the % Fe
determined be affected if the unknowns had not been dried?
2.
For what purpose was sulfuric acid added to the iron samples before titrating?
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